\3Zt 
L 94  3 


THE  EFFECTS  OF  NON-ELECTROLYTES  ON  THE 
SOLVENT  POWER  OF  WATER 


BY 


CLARENCE  B.  LOVELL 


THESIS 


FOR  THE 


Degree  of 


BACHELOR  OF  SCIENCE 


IN 


CHEMICAL  ENGINEERING 


COLLEGE  OF  LIBERAL  ARTS  AND  SCIENCE 
UNIVERSITY  OF  ILLINOIS 


1921 


Digitized  by  the  Internet  Archive 
in  2015 


https://archive.org/details/effectsofnonelecOOIove 


I wish  to  express  my  sincere  thanks  and  appre- 
ciation to  Dr.  J.  H.  Keedy  for  the  interest 
that  he  has  shown  in  thi b investigation,  and 
for  his  cooperation  in  interpreting  the  experi- 
mental data. 


- ‘ 


. - - 


' 


fable  of  Gontenta 


I . 

II  . 
Ill  . 


IV. 


V. 

VI  . 


Introduction  ....... 

Historical  

Experimental 

Materials  ....... 

Procedure  ...... 

Glucose  Solutions  

Glycerol  Solutions  .... 
Acetone  Solutions  . 

Urea  Solutions  ..... 

Di scussion 

Solvent  Action  of  Organic  Solutes 
Group  Influences  ..... 

Hydration  Effects  

Salt  Formation  in  Urea  Solutions  . 
Restrictive  Action  . 


Page 

1 

1 


2 

3 

6 

11 

1? 

30 


31 

94 

34 
37 

35 


Summary 


31 


Bibliography 


Table  of  Plate3. 


Page 


I .  Benzoic  Acid  in  Glucose  Solutions  4 

II . Salicylic  Acid  in  Glucose  Solutions  5 

III.  Specific  Gravity  Curve  of  Glycerol  Solutions  8 

IV.  Benzoic  Acid  in  Glycerol  Solutions  9 

V.  Salicylic  Acid  in  Glycerol  Solutions  10 

VI.  Specific  Gravity  Curve  of  Acetone  Solutions  13 

VII .  Benzoic  Acid  in  Acetone  Solutions  14 

VIII . Salicylic  Acid  in  Acetone  Solutions  15 

IX.  Specific  Gravity  Curve  of  Urea  Solutions  17 

X.  Benzoic  Acid  in  Urea  Solutions  18 

XI.  Salicylic  Acid  in  Urea  Solutions  19 

XII.  Composition  Diagram:  Benzoic  and  Salicylic 

Acids  in  Glucose  Solutions  92 

XIII.  Composition  Diagram:  Benzoic  and  Salicylic 

Acids  in  Glycerol  Solutions  93 

XIV.  Composition  Diagram:  Benzoic  and  Salicylic 

Acids  in  Acetone  Solutions  25 

XV.  Composition  Diagram:  Benzoic  and  Salicylic 

Acids  in  Urea  Solutions  96 


> 
) 

. 


The  Effects  of  h on-Electrolytes  on  the 
Solvent  Power  of  Water . 

I . Introduction* 

The  effects  of  non-electrolytes  on  the  solvent  power  of 
water  have  not  been  developed  very  thoroughly.  That  is  to  say, 
that  while  the  solute  iray  repress  the  solvent  power  of  water, 
the  solute  itself  may  have  a specific  solvent  action,  and  in 
this  way  the  total  solubility  may  be  increased  rather  than 
diminished.  It  is  to  be  expected  that  two  media  which  are 
solvents  in  the  absence  of  each  other  should  retain  their 
solvent  power  when  mixed.  This  specific  solvent  power  of  each, 
however,  may  be  considerably  modified  by  the  presence  of  the 
other.  Furthermore,  solids,  which  in  the  dry  condition  are 
not  considered  solvents,  may  when  dissolved  function  as  sol- 
vents. The  accumulation  of  some  data  on  this  point  has  been 
the  purpose  of  this  investigation. 

II . Historical . 

The  influence  of  dissolved  substances  on  the  solvent 
action  of  water  has  been  extensively  investigated,  both  from 
the  theoretical  and  experimental  standpoints. 

Presuming  the  complete  analogy  between  liquid  and  gaseous 
solutions,  Nernst  developed  the  doctrine  that  "The  concentra- 
tion of  the  molecules  of  the  solute  remains  unchanged  by  the 

(1) 


- 


. 

' 


- 

* 


r . 


p 

presence  of  other  substances ^ The  same  principle  was  as- 
sumed by  A.  A.  Noyes^  in  his  work  on  the  influence  on  solu- 
bility of  the  presence  of  dissociated  substances.  Arrhenius,'3’ 
however,  using  the  data  of  Hojes  on  the  solubility  of  thal- 
lium chloride,  showed  that  this  was  not  even  approximately 
true;  that  the  solvent  power  of  a given  medium  was  diminished 
both  by  the  presence  of  dissolved  molecules  and  by  the  pres- 
ence of  ions.  The  influence  of  ions  in  reducing  solubility 
is  well  illustrated  in  the  familiar  "salting  out"  effect. 

The  theoretical  aspects  of  this  repression  of  solubility  have 
never  been  satisfactorily  developed.  The  attraction  of  the 
electric  charges  of  the  ions  doubtless  bring  about  a contrac- 
tion of  the  solvent  water ( M ©lectrostrietion"4 ) which  would 
cause  a diminution  of  its  solvent  power.  The  withdrawal  of 
the  solvent  water  to  form  hydrates  of  the  molecules  and  ions 
as  suggested  by  Jones®  and  others,  would  result  in  a decrease 
in  the  active  mass  of  the  water,  and  a corresponding  reduction 
in  its  solvent  action.  For  the  most  part,  however,  chemists 
agree  that  a sufficient  explanation  of  this  effect  has  not 
been  given.  Perhaps  all  that  can  be  done  at  the  present  is 
to  accept  the  evasive,  but  helpful  expression  of  Washburn® 
that  the  "thermodynamic  environment"  of  water  is  modified 
by  the  presence  of  dissolved  substances. 

Ill . Experimental . 

Material s --For  this  work,  weak  organic  acids  are  espec- 
ially suitable  because  of  their  rather  low  solubility  in  water, 
their  almost  negligible  ionization,  and  the  ease  by  which 


- 


-7 

• ' 


3 


their  solubility  can  be  accurately  determined  by  titration. 
Benzoic  and  salicylic  acids  were  chosen  as  the  solutes  for 
this  work.  Both  of  these  were  "chemically  pure"  products 
of  the  Merck  brand,  and  their  purity  was  checked  by  the  fol- 
lowing tests:  The  benzoic  acid  melted  at  121° -122° , and  its 

solubility  at  25°  was  found  to  be  0.336  grams  per  100  grams 
of  water.  The  salicylic  acid  melted  at  156°,  and  its  solu- 
bility at  25°  was  found  to  be  0.222  grams  per  100  grams  of 
water.  These  results  agree  well  with  the  accepted  values 
found  in  the  literature.  As  the  non-electrolytes  for  the 
solvent  medium,  glycerol,  glucose,  urea  and  acetone  were 
used.  These  were  of  good  grade,  although  when  dissolved  all 
of  them  except  acetone  showed  a slight  acidity.  Consequently 
blank  titrations  were  run  in  all  determinations  in  these  solu- 
tions, and  corrections  were  made  for  the  acidity  indicated. 
Solutions  of  these  were  made  up  as  needed  in  strengths  vary- 
ing from  1 J to  25^o  by  weight  in  100  parts  of  solution  by 
volume.  Upon  studying  the  curves  obtained,  it  was  found  that 
when  the  results  were  converted  to  a percent-by-vreight  basis 
--  that  is,  grams  per  100  grams  of  solution  --  an  apparently 
simpler  relation  was  brought  out.  The  specific  gravities  of 
these  solutions  were  either  obtained  from  the  "Landolt-Born- 
stein  Tabellen"  or  by  actual  determination,  using  a Westphal 
balance. 

Procedure --These  solutions  were  prepared  in  a closely 
regulated  thermostat.  An  excess  of  the  solid  acid  was  added 
to  200  cc.  of  the  aqueous  solution  in  a 500  cc . Erlenmeyer 
flask.  The  flask  was  stoppered  with  a rubber  stopper  through 


- 


I 

' : e k epnb  fega  4>1* 


■ 


i; 

if 


' 

:: 

:?; 

■*. 


■ 


■ 


- 

i 


i 


4 

;PI<Lte  I 


Solubility  o^B  enzo »c  Acid.  inlUj-ferfent  Percent  Solutions  o\ 

Glucose 


Gm/Per/oo  Gm.  Glucose  Solution 


<p 

09 

o 

■:q 

I 3 

w 

E 

Q) 

cj 

id 

<w 

P- 

25" 


Gnrv Dissolved  Btj  H*0  Be  sent 
in  !00  Gm  Glucose  Solution 


20 


•ntt 


US 


JL 


X 


Om?er/(Jo6m  Solution. 

i 'I  . 1 


.7 


PlfttelT 


Solubility  of  Salicylic  AcidinDlfferent  Percent  Solutions  oj- 

Gtucose. 


20 


to 


GtrnVer  loo& m Glucose  Solution 


0) 

V) 

o 

u 

n 

m 

4t#r 

o 

w 

£ 

H 


GmDissolvedBij  HzO  Present 
in  looGm  G?|u,cose  Solution 


Qm  Per /00Gm  Solution 

I_J5L_I__J I J/. _ I .k_ 


(o 


■3 


!r 


which  a glass  stirrer  was  inserted..  The  stirrer  was  of  the 
centrifugal  type,  revolving  about  200  r.  p.  ip.  The  mixture 
was  stirred  from  10  to  24  hours,  though  it  was  found  that  12 
hours  insured  complete  saturation.  The  excess  of  acid  was  al- 
lowed to  settle  and  25  cc.  portions  were  removed  with  a pipette 
which  was  previously  allowed  to  assume  the  temperature  of  the 
bath.  The  amount  of  acid  dissolved  was  determined  by  titrat- 
ing samples  in  duplicate  with  barium  hydroxide,  using,  phenol- 
phthalein  as  an  indicator.  The  strength  of  the  barium  hydrox- 
ide was  checked  often  against  standard  benzoic  acid  and  no 
appreciable  change  was  noticed.  If  the  excess  of  acid  would 
not  settle  so  as  to  obtain  a clear  sample,  it  was  filtered 
out  by  aspirating  the  liquid  through  a tube  containing  glass 
wool,  according  to  the  method  employed  by  Hoffman  and  Lang- 

7 

beck. 

Glucose  .Solutions- -The  results  for  the  experiments  in 
glucose  solutions  are  shown  in  Table  I , and  the  correspond- 
ing solubility  curves  in  Plates  I and  II . The  black  line  in 

Table  I.  Solubili ties  in  Glucose  Solutions . 

Benzoic  Acid  Salicylic  Acid 


Glucose 
by  Weight 

Solubility 
in  100  g. 
Solution 

Calculated 
Solubility  in 
Water  present 

Solubility 
in  100  g. 
Solution 

Calculated 
Solubility  in 
Water  present 

1o 

G 

G 

G 

G 

0.0 

.3360 

.3360 

.2220 

.2220 

.9975 

.3315 

.3325 

.2180 

.2200 

4.815 

.3205 

.3200 

.2160 

.2110 

9.645 

.3095 

.3030 

.2106 

. 2005 

14.2 

.2980 

.2880 

.2034 

.1905 

18.6 

.28  9 0 

.2730 

.1975 

.1807 

22.8 

.2750 

.2593 

.1915 

.1715 

7 


each  case  was  obtained  by  plotting  per  cent  glucose  against 
the  grams  cf  benzoic  or  salicylic  acid  dissolved  in  100  grams 
of  the  solution.  The  solubilities  represented  by  the  red  line 
are  the  calculated  solubilities  of  the  acid  in  the  amount  of 
water  present  in  100  grains  of  solution,  assuming  of  course 
that  the  glucose  has  no  restrictive  influence  on  the  solvent 
power  of  the  water.  The  red  line  in  all  cases,  then,  is  straight 
since  it  represents  a linear  relation  between  the  weight  of  the 
water  and  the  weight  of  the  acid  dissolved.  The  total-sclubi 1- 
i ty  graph,  as  represented  by  the  black  line,  is  a straight 
line  in  both  cases,  the  solubility  gradually  decreasing  with 
the  concentration  of  the  glucose.  The  difference  in  the  sol- 
ubilities represented  by  the  black  and  red  lines  makes  it  quite 
evident  that  glucose  must  function  as  a solvent.  Hoffman  and 
Langbeck  reported  that  the  presence  of  glucose  in  water  at  25° 
is  without  effect  on  the  solubility  of  benzoic  acid,  but  that 
at  35°  the  solubility  increases  slightly  with  the  concentra- 
tion of  the  glucose.  On  the  other  hand,  they  reported  that  the 
solubility  of  salicylic  acid  increases  regularly  with  the  con- 
centration of  the  glucose,  the  increase  being  the  same  at  35° 
as  at  25°.  Reference  to  Plates  I and  II  shows  that  with  low 
concentrations  of  glucose  the  black  and  red  lines  almost  coin- 
cide, but  diverge  as  the  concentration  of  the  glucose  increases. 
Hoffman  and  Langbeck  worked  with  dilute  solutions  only,  and 
since  the  increase  was  so  slight,  they  probably  assumed  that 
any  deviation  from  a straight  line  was  due  to  experimental 
error,  and  therefore  concluded  that  glucose  was  without  effect, 


‘ 


' /' 


- 


. 

■ 


8 


Plate  HL 


Specific  Gravity  Curve  ForDifferent  Percent  Solutions  of 

Glycerol. 


C) 


o 


<0 


o 


}o^3  3 luaoaad. 


Specific  Gravity  zh 


PlateUT 

Solubility  of  .Benzoic  Acid  inP  iff  e rent  Percent  Solutions  of 

(x\y  cerol. 

144-f-fTTin  Mr  fl  H44*H  44444 1 \ H FtT-I  1 1 1 l+R  r i ITt  4444444414  t4  - M4i  t44-|44  ill  4 f-l  1-44  41444  H44 1 : If  4-rl  i ■ 

I+Sllij -■)  44444  :4  ~4— 


1 (*m  Per/0O(*m  Glycerol Solution 

— L_|Gm  Dissolved  By  HaO  Present  in 

£ ||2PGm&\i)CGrol  Solution 

05 

■ I 


(U 


10 

PlotelT 


5o)iibilitij  oj-  Salicy)  ic  /4cid.  in  Different  Percent  Solutions  oj- 

&\yce rol 


Q> 

O 

3"i 

CS 

;-M 

L 

W 

u 

r-WT 

* 

Z5 


Gim  Per IOOGcrr\  (*lijceiro\  So\u.tion 

GirnDissoWed-By  H*0  Present  in 
/oo&m  GiUjcerot  Solution 


20 


JS 


10 


Gtrn  YerlQQ&m  Solittipa; 

-J #_1 


I 


11 


especially  in  the  case  of  benzoic  acid.  Another  interesting, 
fact  is  that  benzoic  is  more  soluble  than  salicylic  acid. 

This  was  the  case  in  all  determinations  and  is  in  accord  with 
Bourgoin’s®  work.  He  found  that  below  40°  salicylic  acid  is 
less  soluble  than  benzoic,  while  above  40°  it  is  more  soluble. 

Glycerol  Solutions--The  results  of  solubility  determi- 
nations in  glycerol  solutions  are  shown  in  Table  II , and  the 
corresponding  solubility  curves  in  Plates  IV  and  V.  These 
data,  if  plotted  on  the  grams-per-lOG  cc.  basis,  give  a curved 


Table  II . Solubilities  in  Glycerol  Solutions . 

Benzoic  Acid  * Salicylic  Acid 


Glycerol 
by  Weight 

Solubility 
in  100  g. 

Calculated 
Solubility  in 

Solubility 
in  100  g. 

Calculated 
Solubility  in 

Solution 

Water  present 

Solution 

Water  present 

at 

/° 

G 

G 

G 

G 

0.0 

.3360 

.3360 

.2220 

.2220 

.998 

.3365 

.3330 

. 2225 

.2195 

4.95 

. 3450 

.3200 

.2260 

.2120 

9.765 

.3520 

.3030 

.2340 

.2000 

14.5 

.36  20 

.2875 

.2410 

.1898 

19.1 

.3730 

.27  20 

.2495 

.1795 

S3. 5 

.3839 

.2570 

.8565 

.1698 

line.  On 

the  other  hand,  the  graph  obtained  on 

a basis  of 

grams  per 

ICO  grams  of 

solution  is  a 

straight  line.  The  sol- 

ubility  cf 

benzoic  acid 

and  salicylic 

acids  in  glycerol  was 

determined 

in  the  usual 

way,  and  values  of  1.93 

grams  and 

1.59  grams,  respectively,  per  100  grams  of  solvent.  From  these 
values  the  solubilities  of  these  acids  may  be  calculated  for 
the  glycerol  fraction  of  the  solution,  just  as  it  was  calcti- 
lated  for  the  water  fraction,  on  the  aseiimption  that  there  is 


. 


' r.-  9 


. 


1? 

# 

no  restrictive  influence  involved.  For  example,  in  100  grams 
of  a 23.5 fc  solution  of  glycerol,  the  solubility  of  benzoic 
acid  in  the  water  fraction  (76.5  grams)  would  be  .2570  grams; 
in  the  glycerol  fraction  (23.5  grams)  .4536  grains.  The  total 
solubility  calculated  on  the  assumption  of  no  restrictive  in- 

f 

fluence  would  then  be  .7365  grains.  The  value  actually  found 
is  .3839  gram.  The  difference,  .3267  grams,  represents  the 
total  repression  of  solubility  in  this  solution.  In  the  same 
way  the  restrictive  influence  for  salicylic  acid  in  100  grams 
of  23.5$  glycerol  solution  was  calculated  to  be  .2870  grams. 
This  conclusively  shows  that  solubility  is  not  an  additive 
quantity;  that  while  both  components  of  the  solution  exert  a 
solvent  action,  there  is  seme  restrictive  influence  in  opera- 
tion. The  solubility  curves  in  glycerol  and  glucose  solutions 
are  very  similar,  which  strongly  suggests  that  analogous  rela- 
tionships exist  in  both  cases.  The  striking  difference  is, 
glycerol  solutions  are  the  better  solvents. 

Acetone  Solutions--The  results  of  the  solubility  deter- 
minations in  acetone  solutions  are  given  in  Table  III,  and 
the  corresponding  solubility  graphs  in  Plates  VII  and  VIII. 

The  solubilities  increase  very  rapidly  for  the  higher  concen- 
trations of  acetone.  As  in  the  case  of  glycerol  solutions 
the  total  solubility  is  not  the  sum  of  the  solubilities  of 
the  water  and  acetone  fractions,  considered  separately. 

Upon  plotting  the  solubilities  on  a ” composition”  diagram 
(See  Plate  XIV,  page  25),  curves  resulted  which  are  similar 


- 


, 


* % 


13 


ior\$ 


j±i±i 


14 


Grm  Per /oo  Gem  Solution. 


15 

PlateMT 


Solubility  of  Saiicijlic  Acid,  in  Different  Percent  Solutions  of 

Acetone 


vO 


»n 


(*m  Per  ioo(krr\  of  Solution 


17*k- 


. GrnVzr  100 Urea  So  tut  von 


|SolabiUtij  of  Benzoic  Acid  inDif  jerervll  Percent  Solutions  of 

llfea. 


frrH— 


« 


\gr 


t*> 


C* 


I 


'oajiyi  }uao3^ 


^1 


$L 


m 


Gm  Ver/oo(km  of  Solution. 


(* mper  lOO&m  Urea  Solution 


19 

piatcxi 


Solubilituof  SaUcuhc  Acid,  in  Different  Percent  Solutions  of 

Llrea. 


£ 

d 

(n 

Q) 

ft  c 
o 2 

•<  4-* 

m 3 


CO 


T>3yrn 


*o 


&m  Per looGm  of  Solution 


.. 


; 


Table  III.  Solubilities  in  Acetone  Solutions. 


20 


Benzoic  Acid  Salicylic  Acid 


Acetone 
by  Weight 

Solubi li ty 

Calculated 

Solubility 

Calculated 

in  100  g. 

Solubility  in 

in  106  g. 

Solubility  in 

Solution 

Water  present 

Solution 

Water  present 

* 

G 

G 

G 

G 

0.0 

.3360 

.3360 

.2820 

.8220 

1 .0025 

.3610 

.3329 

.2320 

.2200 

5.04 

.5150 

.3190 

.3320 

.2110 

10 . 18 

.8  850 

.3020 

.5510 

.1995 

15.35 

1.3900 

.2846 

.9220 

.1880 

80.65 

2.4650 

.2663 

1.6100 

.1760 

86.00 

5.1200 

.2489 

3.4600 

.1644 

to  those  obtained  by  Seidell9  for  ethyl  alcohol  solutions. 

He,  however,  offered  no  explanation  why  these  graphs  assumed 
such  a shape.  It  is  also  to  be  noticed  that  the  curves  of 
both  acids  when  drawn  on  a basis  of  grains  of  acid  per  IOC 
grams  of  solution  have  practically  the  same  slope. 

Urea  Solutions — The  results  of  solubility  determinations 
in  urea  solutions  are  shewn  in  Table  IV,  and  the  solubility 
graphs  in  Plates  X and  XI.  The  solubility  of  both  acids  in- 
creases rather  rapidly  with  concentration  of  urea.  Both 

Table  IV.  Solubilities  in  Urea  Solutions . 


Benzoic  Acid  Salicylic  Acid 


Solubility 

Calculated 

Solubility 

Calculated 

urea 

in  100  g. 

Solubility  in 

in  100  g. 

Solubility  in 

by  Weight 

Solution 

Water  present 

Solution 

Water  present 

% 

G 

G 

G 

G 

0.0 

.3360 

.3360 

.2220 

.2220 

.997 

.3455 

.3330 

. 8290 

.2200 

4.93 

.4040 

.3195 

.3043 

.2110 

9.75 

.4790 

.3030 

.3860 

.2005 

14.45 

.57  25 

.2880 

.4710 

.1900 

19.00 

.6700 

.2720 

.4775 

.1800 

23.48 

.7650 

. 257  2 

♦ 48  25 

.1700 

* 

• 

* 

. • 

. 

. 

V 

• 

* 

. 

. 

- 

■* 

. ‘ 

or' 

. ' 

. 


. 


. 


curves  bend  slightly  up  to  a urea  concentration  of  about  14.45??, 
and  then  an  inflection  takes  place,  not  very  marked  in  the  ben- 
zoic acid  curve,  but  sharp  in  the  salicylic  acid  curve.  The 
peculiarities  of  these  graphs  will  be  discussed  later. 

I V . Discussj on . 

Solvent  Action  of  Organic  Solutes --From  the  solubilities 
as  determined  above,  it  is  seen  that  the  solubilities  of  ben- 
zoic and  salicylic  acids  in  a definite  amount  of  solution  is 
decreased  by  the  presence  of  glucose,  but  increased  by  the  pres- 
ence of  glycerol,  urea  and  acetone.  The  concentration  of  the 
water  in  the  glucose  solutions,  however,  is  greatly  reduced. 

For  example,  100  cc.  of  a 25%  glucose  solution  contains  only 
84.46  grains  of  water.  From  this  it  may  be  seen  that  the  sol- 
ubility in  glucose  solutions  is  too  great  to  be  accounted  for 
by  the  solvent  action  of  the  water  actually  present.  It  fol- 
lows, therefore,  that  either  the  presence  of  the  glucose  in- 
creases the  solvent  power  of  water,  or  that  glucose  itself 
has  a solvent  action  on  the  acids.  The  first  assumption  is 
rejected  since  it  is  at  variance  with  the  work  of  Noyes°  and 
Arrhenius^,  who  found  that  the  presence  of  dissolved  substances, 
especially  of  ions,  tend  to  reduce  rather  than  increase  the 
solvent  power  of  water.  On  the  other  hand,  it  is  not  sur- 
prising that  an  organic  substance  like  glucose  should  be  a 
solvent  for  benzoic  and  salicylic  acids.  This  viewpoint  is 
supported  by  the  following  experiment:  Some  solid  glucose 

and  benzoic  acid  were  carefully  heated  together  in  a test 


. : 


, 


•- 


. 


- 

- » i 


■ 

■ 

HtO  ioo  °fo  Percent  Composition  Glucose  m% 

10  X'O  *\Q  j\Q  fo  L\0  1 i jo  . ' 60  . j ft?  ..  i . . 1 too 


Hxo  too*/*  Percent  Composition  G\\\cero\  iqqJo 


°4 


tube,  giving  a perfectly  homogeneous  mixture.  There  was  a 
slight  darkening  during  the  process  due  to  the  formation  of 
caramel  and  a little  steam  was  evolved.  Upon  cooling  a layer 
of  benzoic  acid  separated  on  top  of  the  melt. 

Group  Inf luences- -Certain  inferences  may  be  drawn  as  to 
the  effect  of  certain  groups  on  the  solvent  power  of  an  or- 
ganic solvent  for  benzoic  and  salicylic  acids.  First  of  all, 
it  should  be  noticed  that  the  presence  of  OH-groups  reduces 
the  solubility  of  these  acids,  as  evidenced  by  their  low  sol- 
ubility in  glucose  and  glycerol  solutions.  The  fact  that 
glucose  is  a poorer  solvent  than  glycerol  is  doubtless  due 
to  the  larger  number  of  OH-groups  in  the  former.  The  increased 
solubility  in  acetone  solutions  might  be  anticipated  from  the 
two  CH^-groups,  since  both  acids  are  fairly  soluble  in  hydro- 
carbons. The  CO-group  is  probably  hydrated  in  solution,  be- 
coming C(GH)p,  and  has  a repressing  influence  on  solubility. 

As  100%  acetone  is  approached,  however,  (see  Plate  XIV,  page 
°5 ) , the  solubility  becomes  greater,  indicating  that  anhydrous 
acetone  is  a much  better  solvent  than  the  hydrate.  This  leads 
to  the  conclusion  that  the  CO-group  has  a greater  solvent  ac- 
tion on  these  acids  than  the  C(OK)s>  -group.  From  the  same 
point  of  view,  the  high  solubility  in  urea  solutions  is  to 
be  attributed  to  the  NH^-groups.  As  confirmatory  evidence 
of  this  may  be  cited  the  fact  that  aniline,  CgHg.NH^,  is  a 
better  solvent  for  these  acids  than  benzene,  CgHg. 

Hydration  Effects- -Jones  and  his  collaborators  have  in- 


J 

. 


. L. 

' 


- 


Composition  Gurve  ofB  enzoic  and.  Sal  icy  Uc  Actd-*  in  Acetone  Solution 


Composition  Curve  of  Benzoic  and  Salicijic  Acids  in  lire  a Solutions 


H*0  /ooy0  Percent  Composition  Urea  my0 


sis ted  strongly  that  the  deviation  of  solutions  from  what  might 
be  called  their  "ideal"  behavior  is  due  to  hydration  effects. 
According  to  this  "solvation”  hypothesis,  molecules  of  the 
solvent  combine  with  solute  molecules  and  ions  to  form  loose 
molecular  compounds  called  "solvates”.  This  theory  gives  the 
most  immediate  explanation  of  the  fact  that  the  solvent  power 
of  solutions  does  not  follow  the  law  of  mixtures.  The  water 
is  partly  consumed  in  forming,  hydrates  of  the  solute,  thereby 
reducing  the  number  of  effective  molecules.  In  many  cases  it 
is  probable  that  a series  of  hydrates  is  formed,  each  having 
a specific  solvent  action  of  its  own.  To  illustrate:  Sei- 

dell9 found  that  camphoric  acid  in  water-ethyl  alcohol  mixtures 
had  a maximum  solubility  in  about  87$  alcohol.  This  indicates 
that  alcohol  molecules  of  a certain  degree  of  hydration  have 
a higher  solvent  action  than  the  pure  alcohol.  For  the  other 
acids  whose  solubility  he  measured  in  water-alcohol  mixtures, 
the  maximum  solubility  was  for  100$  alcohol.  This  change  in 
hydration  makes  solubility  effects  all  the  more  complex,  and 
since  probably  other  influences  are  involved,  it  may  be  suf- 
ficient to  attribute  this  deviation  from  ideal  behavior  to  a 
modified  thermodynamic  environment  of  the  solution. 

Salt  Formation  in  Urea  Solutions — The  possibility  of  salt 
formation  in  the  case  of  urea  solutions  is  very  strong,  and  it 
is  altogether  likely  that  the  increasing  solubility  of  the 
acids  with  increase  in  the  urea  concentration  is  largely  due 
to  this  effect.  Such  action  would  affect  the  solubility  of 


°8 

the  acids  in  two  opposing  ways,  viz.:  The  withdrawal  of  urea 

from  the  solution  to  form  salts  would  reduce  the  solvent  power 
of  the  solution,  since  urea  seems  to  have  a high  specific  sol- 
ubility for  these  substances;  on  the  contrary,  the  withdrawal 
of  the  acid  by  salt  formation  would  increase  its  solubility. 
Attempts  to  prepare  urea  benzoate  and  urea  salicylate  failed. 
h either  of  them  has  been  reported  in  chemical  literature,  and 
so  their  existence,  it  must  be  granted,  is  problemmatical . If 
they  do  exist  in  solution  they  are  probably  rather  soluble  and 
considerably  hydrolyzed  by  the  water.  Upon  attempting  * to  re- 
move the  water  by  boiling,  urea  only  would  remain  since  both 
of  the  acids  are  readily  volatile  in  steam;  if  the  water  were 
removed  by  evaporation  at  low  temperatures,  the  free  acids, 
by  reason  of  their  low  solubilities,  would  tend  to  separate 
first . 

The  rather  abrupt  break  in  the  solubility  curve  of  sal- 
icylic acid  in  urea  solutions  (see  Plate  XV,  page  °6)  may  be 
easily  interpreted  by  means  of  the  Fhase  Rule  in  the  light  of 
the  assumption  of  salt  formation.  In  a urea  solution  contain- 
ing an  excess  of  salicylic  acid,  there  are  (counting  the  water) 
three  components.  The  phases  are  three  — ■ solid  salicylic 
acid,  solution,  vapor.  Therefore,  substituting  in  the  ex- 
pression for  the  Phase  Rule,  F = G + 2 - P,  we  obtain  F = 2. 
That  is,  there  are  two  degrees  of  freedom.  From  this  it  would 
follow  that  for  a definite  temperature,  25°,  there  should  be 
a definite  solubility  of  salicylic  acid  for  every  concentra- 
tion of  urea.  But,  by  reference  to  the  figure,  it  will  be 


i 

. 

* 


_ 


— 

- 


seen  that  above  14.45^urea  concentration  the  solubility  of  the 
acid  becomes  practically  constant.  This  means  that  another 
phase  has  been  added  to  the  system,  presumably  solid  urea 
salicylate,  so  that  now  F = 1. 

There  is  to  be  mentioned  another  complication  in  dealing 
with  urea  solutions.  The  decomposition  of  urea  into  ammonia 
and  carbon  dioxide  is  catalyzed  by  acids,  the  latter  uniting 
with  the  ammonia  to  form  ammonium  salts.  That  this  does  occur 
was  shown  qualitatively  by  adding  an  excess  of  salicylic  acid 
to  100  cc.  of  25%  urea  solution.  Upon  standing  for  several 
days  at  temperatures  slightly  above  that  of  the  room  the 
excess  of  acid  gradually  went  into  solution  with  an  evolution 
of  carbon  dioxide.  While  ammonium  salicylate  is  almost  neu- 
tral and  would  affect  the  titration  results  only  slightly, 
it  may  considerably  influence  the  solubility  of  the  molecular 
salicylic  acid. 

Restri cti ve  Action- -The  simplest  mode  of  behavior  when 
a substance  is  dissolved  in  a mixture  of  two  mutual  solvents 
is  that  each  of  the  solvents  should  retain  its  specific  sol- 
vent power.  In  such  a case  the  solubility  of  the  substance 
should  be  additive  and  would  follow  the  law  of  mixtures.  The 
solubility  graph  in  such  a case  would  be  represented  by  a 
straight  line,  as  shown  diagrammatically  in  Figure  1 (see  next 
page),  where  A and  B are  the  two  solvents  and  the  solubility 
of  the  solute  C is  represented  by  the  ordinates.  The  straight 
line  DE  would  represent  this  ideal  behavior  — that  is,  when 


. - 


. 


■ 


- 

- 

» 

30 


too?,  ft 


/oo%B 


the  solvent  power  of  one  solute  is 
not  changed  by  the  presence  of  the 
other.  More  generally  the  solubility 
graph  is  a curve  of  the  general  form 
DFE,  as  will  be  discussed  later.  Ben- 
zoic acid  and  salicylic  acid  in  glu- 
cose solutions  (see  Plate  XII,  page 
pP)  give  graphs  that  appear  straight 
lines  up  to  a concentration  of  50^ 
glucose,  which  is  somewhat  above  the  limiting  concentration 
for  satisfactory  work.  This  may  be  due  to  the  fact  that  the 
curvature  is  inappreciable,  and  too  much  significance  should 
not  be  given  to  the  solubility  values  for  pure  glucose  as  in- 
dicated by  the  projection  of  these  graphs.  More  probably  the 
true  values  would  be  considerably  greater.  On  account  of  glu- 
cose being  a solid  at  25°  no  attempt  was  made  to  determine 
these  values  directly. 


Figure  1 


In  the  case  of  binary  solvents,  the  general  condition  is 
that  solubility  is  not  an  additive  quantity,  and  does  not  fol- 
low the  law  of  mixtures.  Each  solvent  seems  to  restrain  the 
solvent  action  of  the  other,  so  that  the  combined  solubility 
is  less  than  if  each  were  operating  separately.  The  graph  is 
not  a straight  line  but  a curve,  of  the  general  shape  DFE. 


An  expression  for  the  solubility  of  a substance  in  a bi- 
nary system  of  solvents  may  be  developed  as  follows:  Let  A 

and  B be  the  two  solvents,  where  a and  b represent  the  number 


. j 

■ 


of  grans  (or  moles)  of  each  present,  and  3 the  solubility. 

The  expression  for  the  total  solubility  is  then 

S = aS^  + bS-g  - abk. 

The  last  term  represents  the  repression  of  solubility  and  has 
not  justification  beyond  the  assumption  that  it  is  dependent 
on  the  concentrations  of  both  A and  B.  This  last  term  becomes 
nil  whenever  a or  b is  zero  — that  i 3,  whenever  the  solute 
is  a single  substance.  The  graph  of  the  above  equation  is 
of  the  general  shape  of  the  curve  DPS  in  Figure  1.  By  using 
the  data  for  a given  curve  values  for  k may  be  derived  which 
are  roughly  constant.  Table  V shows  the  values  calculated 
from  the  foregoing  solubility  data. 


Table  V.  Values  for  k in  £>  = aSA  + bSp  - abk. 


Benzoic 

Acid 

Salicylic 

Acid 

entration 

in 

in 

in 

in 

Solution  1 

Glycerol 

Acetone 

Glycerol  Acetone 

1%  k 

= .00243 

k = .00273 

k = .00202  k 

= .00  281 

5 jo 

.00253 

.00276 

.00208 

.00301 

10  fc 

.00264 

.00279 

.00217 

.00309 

155$ 

.00  274 

.00279 

.00226 

.00314 

20  $ 

♦ 00  286 

.00250 

.00236 

.00308 

.00300 

.00157 

.00247 

.00254 

Average 

, .00  270 

.00252 

.00226 

.00  295 

V . Summary . 

1.  The  solubilities  of  benzoic  and  salicylic  acids  have 
been  measured  in  solutions  of  glucose,  glycerol,  acetone  and 
urea  at  25°  for  the  range  of  ifo  to  23.5^6  by  weight. 

2.  These  organic  solutes  exert  a solvent  action  on  ben- 


* 

- - 

' 


32 

zoic  and  salicylic  acids. 

3.  The  influence  of  certain  groups  on  solubility  has  been 
considered. 

4.  The  total  solubility  in  all  cases  was  found  to  be 
less  than  the  combined  solubilities  of  the  components  operating 
separately . 

5.  This  repression  of  solvent  action  i3  probably  largely 
due  to  hydbation  effects. 

6.  The  formation  of  urea  salicylate  in  urea  solutions, 
as  interpreted  by  the  Phase  liule,  was  indicated. 

7.  An  empirical  expression  for  solubility  in  a binary 
medium  has  been  developed. 


' 

' 


. 


33 


VI  . Bibliography . 

1.  Hernst:  Theoretical  Chemistry,  p.  357  (1911). 

2.  A.  A.  Hoyes:  Z.  physik.  Chem.,  6,  243  (1890). 

3.  Arrhenius:  Z.  physik.  Chem.,  51,  824  (1899). 

4.  Drude  und  Herns t:  Z.  physik.  Chem.,  P5,  79  (1894). 

5.  H.  C.  Jones:  Z.  physik.  Chem.,  15,  419  (1894). 

6.  Washburn:  Principles  of  Physical  Chemistry,  p.  179  (1915). 

7.  Hoffman  und  Langbeck:  Z.  physik.  Chem.,  51,  385-434  (1905)* 

8.  Bourgoin:  Compt.  rend.,  87,  62-64  (1903). 

9.  Seidell:  Trans.  Am.  ■Electroch.  Soc.,  13,  319-331  (1908). 


